Molar Mass Calculator
Calculate the molar mass of compounds based on their chemical formula
Supports element symbols (H, O, Na, Ca, etc.) and number subscripts, e.g., H2O, CaCO3
Quick Input for Common Compounds:
Frequently Asked Questions
How to write chemical formulas correctly?
Rules for writing chemical formulas: (1) Element symbols start with uppercase letter, second letter lowercase (e.g., Na, Ca, Fe). (2) Numbers written as subscripts after element symbols indicate atom count (e.g., H₂O, CO₂). (3) Groups in parentheses with multiple atoms have numbers outside parentheses (e.g., Ca(OH)₂, not yet supported by this calculator). (4) Organic compounds written conventionally (e.g., CH₃COOH for acetic acid). Note: This calculator currently does not support parentheses, they need to be expanded. For example, Ca(OH)₂ should be written as CaO2H2.
What is the difference between molar mass and molecular weight?
Molar mass and molecular weight are closely related but have subtle differences: (1) Molar Mass: Unit is g/mol, is the mass of 1 mole of substance. (2) Molecular Weight: Unitless (relative quantity), is the relative mass of one molecule. Numerically, molar mass and molecular weight are equal, but the units differ. For example, the molecular weight of water is about 18, and the molar mass is about 18 g/mol. In practical applications, these two terms are often used interchangeably because their numerical values are the same.
What is Avogadro's constant?
Avogadro's number (NA) refers to the number of basic units contained in 1 mole of substance: NA = 6.02214076 × 10²³ mol⁻¹. Significance: (1) 1 mole of carbon-12 atoms contains approximately 6.022 × 10²³ atoms, with a mass of exactly 12 grams. (2) This is the bridge connecting microscopic particle counts to macroscopic mass. (3) Makes stoichiometric calculations possible. This constant is named after Italian scientist Amedeo Avogadro, who proposed in 1811 that equal volumes of gases at the same conditions contain equal numbers of molecules.
How are atomic masses determined?
Atomic mass (relative atomic mass) is defined based on the carbon-12 isotope: (1) Carbon-12 atomic mass is defined as exactly 12. (2) Other elements' atomic masses are measured relative to carbon-12. (3) Takes into account the natural abundance-weighted average of all existing isotopes. (4) Is a unitless relative value. Why are atomic masses usually not integers? Because most elements have multiple isotopes (same proton number, different neutron numbers), the atomic mass is the weighted average of the masses of all isotopes. For example, chlorine has two main isotopes Cl-35 and Cl-37, with an atomic mass of 35.45.
How to use molar mass in calculations?
Molar mass has wide applications in chemical calculations: (1) Mass and mole conversion: Mass = moles × molar mass. Example: Mass of 2 moles of water = 2 mol × 18.016 g/mol = 36.032 g. (2) Solution preparation: Mass = molarity × volume × molar mass. Example: To prepare 1L of 0.1M NaCl solution, need NaCl = 0.1 mol/L × 1 L × 58.44 g/mol = 5.844 g. (3) Chemical reaction stoichiometry: Calculate mass relationships between reactants and products based on balanced chemical equations and molar masses.
What are the practical applications of molar mass?
Molar mass has extensive practical applications: (1) Stoichiometric calculations: Calculate mass relationships between reactants and products in chemical reactions. (2) Solution preparation: Use solute molar mass to calculate required mass when preparing solutions by molarity. (3) Drug dosage calculations: Calculate drug doses based on molar mass in medicine, especially for active ingredient amounts. (4) Biochemistry: Calculate molecular weights of biomolecules like proteins and DNA to study their structure and function. (5) Materials science: Molecular weight characterization of polymers and materials affects their physical and chemical properties. (6) Analytical chemistry: Calculate content of components in samples in quantitative analysis.
Disclaimer
The molar mass provided by this calculator is based on standard atomic weight data and is for reference only. For complex compounds or calculations with special requirements, please consult professional chemical literature or databases. For medical or pharmaceutical applications, please verify the calculation results.
Molar mass converts between the mass of a substance you can weigh and the number of particles it contains, which is the bridge every quantitative calculation in chemistry crosses. It is numerically equal to the relative molecular mass but carries units of grams per mole, and that unit distinction is the source of most of the confusion around it.
The formula
Molar mass (g/mol) = sum over all atoms of (atomic weight x count)
Moles = mass (g) / molar mass (g/mol)
Particles = moles x 6.02214076 x 10^23- Atomic weights are the standard values published by IUPAC, which are averages weighted by natural isotopic abundance
- The Avogadro constant is exact by definition since the 2019 SI redefinition
Atomic weights: IUPAC Commission on Isotopic Abundances and Atomic Weights. The mole is defined as exactly 6.02214076 x 10^23 elementary entities.
How the calculation works
The molar mass of a compound is the sum of the atomic weights of its constituent atoms, each multiplied by how many times it appears in the formula. For water, H2O, that is two hydrogens at 1.008 plus one oxygen at 15.999, giving 18.015 g/mol. The arithmetic is simple; the errors almost always come from parsing the formula rather than from the addition.
Parentheses and hydrates are where mistakes cluster. In calcium nitrate, Ca(NO3)2, the subscript 2 applies to everything inside the brackets, so there are two nitrogens and six oxygens. In a hydrate such as CuSO4 5H2O, the water molecules are part of the formula unit and must be included, which adds about 90 g/mol.
| Compound | Formula | Molar mass (g/mol) |
|---|---|---|
| Water | H2O | 18.015 |
| Sodium chloride | NaCl | 58.44 |
| Glucose | C6H12O6 | 180.156 |
| Sulfuric acid | H2SO4 | 98.079 |
| Calcium carbonate | CaCO3 | 100.087 |
| Copper(II) sulfate pentahydrate | CuSO4 5H2O | 249.685 |
Why atomic weights are not whole numbers
Most elements exist as a mixture of isotopes with different neutron counts, and the standard atomic weight is the abundance-weighted average across that mixture as found naturally on Earth. Chlorine is 35.45 rather than 35 because natural chlorine is roughly three parts chlorine-35 to one part chlorine-37.
This is why a molar mass calculation and a monoisotopic mass calculation give different answers, which matters in mass spectrometry. A mass spectrometer resolves individual isotopes and reports the mass of a specific isotopic composition; a balance weighs the natural mixture. Using one where the other is required is a common source of error in analytical work.
- IUPAC publishes some atomic weights as intervals rather than single values, because natural isotopic composition varies measurably by source.
- For elements with no stable isotope, the value quoted is conventionally that of the longest-lived isotope.
What molar mass is used for
The two everyday uses are converting a weighed mass to moles for a reaction, and preparing a solution of known concentration. To make one litre of 0.1 molar sodium chloride, multiply 0.1 mol by 58.44 g/mol to get 5.844 g. Every stoichiometric calculation runs through the same conversion.
In clinical and pharmaceutical contexts, molar mass is what converts between mass concentration and molar concentration, which is why the same laboratory result can be reported as mg/dL in one country and mmol/L in another. The conversion factor between those two units for any analyte is derived directly from its molar mass.
Limitations
- Standard atomic weights are averages over natural isotopic abundance. Materials with non-natural isotopic composition, including deuterated or enriched compounds, require different values.
- Molar mass is not the same as monoisotopic mass. Mass spectrometry uses the latter, and the difference grows with molecule size.
- The calculator depends on the formula being entered correctly. Ambiguous or incorrectly bracketed formulas produce confidently wrong answers.
- Hydrates must include their water of crystallisation, which is easy to omit and can change the result by a third or more.
- For polymers and other substances without a single defined formula unit, the concept of a single molar mass does not apply in the same way; average molar masses are used instead.
- IUPAC revises standard atomic weights periodically, so values from different eras differ slightly in the final digits.
Frequently asked questions
How do I calculate molar mass?
Add the atomic weight of every atom in the formula, multiplied by how many times it appears. Water, H2O, is two hydrogens at 1.008 plus one oxygen at 15.999, giving 18.015 g/mol.
What is the difference between molar mass and molecular weight?
They are numerically the same but conceptually different. Molecular weight is a dimensionless ratio relative to one twelfth of a carbon-12 atom; molar mass carries units of grams per mole and applies to a mole of substance.
Why is chlorine 35.45 and not 35?
Because standard atomic weights are averages weighted by natural isotopic abundance. Natural chlorine is roughly three parts chlorine-35 to one part chlorine-37, and the average of that mixture is 35.45.
How do I handle a formula with parentheses?
The subscript after a closing bracket multiplies everything inside it. In Ca(NO3)2 there are one calcium, two nitrogens and six oxygens, because the 2 applies to both the N and the O3.
Do I include the water in a hydrate?
Yes. In CuSO4 5H2O the five water molecules are part of the formula unit and add about 90 g/mol. Omitting them is one of the most common errors in molar mass calculations.
References
- International Union of Pure and Applied Chemistry. Periodic Table of Elements and standard atomic weights.
- National Institute of Standards and Technology. Periodic Table of the Elements.
- National Institute of Standards and Technology. Atomic Weights and Isotopic Compositions.
- NIST. The International System of Units (SI), definition of the mole.